A base without OH⁻
In lesson 1.8.4 you met common acids and bases: acids turn litmus red, and bases turn phenolphthalein pink. Lesson 2.7.5 defined acids and bases by the ions they release in water. But a few facts don’t fit that definition well:
- Ammonia, NH₃, contains no OH⁻, yet its solution is basic.
- Sodium carbonate, Na₂CO₃ (sold as washing soda), is a salt, yet its solution turns phenolphthalein pink.
- When open bottles of concentrated hydrochloric acid and ammonia are placed side by side, a white smoke of NH₄Cl forms where the HCl and NH₃ gases meet. There is no water, yet it looks a lot like neutralization.
This lesson reviews the Arrhenius definition, then introduces a broader one.
The Arrhenius definition
In 1887 the Swedish chemist Svante Arrhenius proposed the Arrhenius definition电离理论, the one used in lesson 2.7.5:
- An acid is a compound that produces H⁺ ions in water, for example HCl → H⁺ + Cl⁻.
- A base is a compound that produces OH⁻ ions in water, for example NaOH → Na⁺ + OH⁻.
In this picture, neutralization is H⁺ + OH⁻ → H₂O. The Arrhenius definition is simple and useful, but it only covers water solutions and only counts compounds that release H⁺ or OH⁻, so it can’t explain the examples above.
The Brønsted–Lowry definition
In 1923 Johannes Brønsted in Denmark and Thomas Lowry in England independently proposed the Brønsted–Lowry definition质子理论:
- An acid is a proton donor: it gives up H⁺.
- A base is a proton acceptor: it takes H⁺.
Why is H⁺ called a proton? A hydrogen atom is one proton and one electron; take away the electron and only a proton is left. In water, H⁺ never exists on its own. It attaches to a water molecule to form the hydronium ion水合氢离子, H₃O⁺. So when HCl dissolves in water, HCl actually hands H⁺ to H₂O:
HCl donates H⁺, so it is the acid; H₂O accepts H⁺, so it is the base. HCl → H⁺ + Cl⁻ is shorthand for this reaction.
A Brønsted–Lowry acid–base reaction is a proton transfer from the acid to the base. No water is needed, and ions can be acids or bases too. The white smoke above forms when HCl gives H⁺ to NH₃.
Conjugate acid–base pairs
Look at what happens when ammonia dissolves in water:
H₂O gives up H⁺, so it is the acid; NH₃ takes H⁺, so it is the base. The reaction is reversible: in the reverse reaction NH₄⁺ gives H⁺ back to OH⁻, so NH₄⁺ is also an acid and OH⁻ is also a base.
- The species a base becomes after it accepts H⁺ is its conjugate acid共轭酸: NH₄⁺ is the conjugate acid of NH₃.
- The species an acid becomes after it donates H⁺ is its conjugate base共轭碱: OH⁻ is the conjugate base of H₂O.
An acid and its conjugate base form a conjugate acid–base pair共轭酸碱对. The two members differ by exactly one H⁺. Every proton-transfer reaction has two pairs; here they are NH₄⁺/NH₃ and H₂O/OH⁻.
Now the other puzzles make sense. NH₃ accepts H⁺, so it is a base without containing OH⁻. CO₃²⁻ takes H⁺ from water (CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻) and leaves OH⁻ behind, so sodium carbonate solution is basic.
Species that can go either way
The same species can donate H⁺ to a stronger base and accept H⁺ from a stronger acid. In the two reactions above, water is the base in the first and the acid in the second. HCO₃⁻ behaves the same way: with OH⁻ it gives up H⁺ and becomes CO₃²⁻, and with H₃O⁺ it accepts H⁺ and becomes H₂CO₃. A species that can either donate or accept a proton is amphiprotic两性物质.
So to decide whether a species is an acid or a base, look at whether it gives or takes H⁺ in that reaction.
Here are seven reactions. Tap each species to label it acid, base, conjugate acid, or conjugate base, then press “Check” to see where the H⁺ goes:
How Chinese textbooks differ
- Chinese textbooks mostly use the Arrhenius (ionization) definition and write H⁺ in equations; the Brønsted–Lowry definition is extension material. US textbooks and AP mainly use Brønsted–Lowry and often write H₃O⁺.
- Chinese textbooks write aqueous ammonia as NH₃·H₂O and its ionization as NH₃·H₂O ⇌ NH₄⁺ + OH⁻.
- In Chinese textbooks, “两性” (amphoteric) mostly refers to substances like Al₂O₃ and Al(OH)₃ that react with both acids and strong bases. US textbooks use both amphiprotic and amphoteric for species like H₂O and HCO₃⁻.
- There is an even broader definition, the Lewis definition, based on electron pairs. It isn’t covered here.
Example 1: one reaction, two definitions
Hydrogen chloride and ammonia gases react to form a white smoke: NH₃(g) + HCl(g) → NH₄Cl(s). Is this an acid–base reaction under the Arrhenius definition? Under the Brønsted–Lowry definition?
- The two definitions. Arrhenius: an acid produces H⁺ in water and a base produces OH⁻ in water. Brønsted–Lowry: an acid donates a proton (H⁺) and a base accepts one; no water is required.
- Arrhenius. The reaction happens between gases, with no water. Neither HCl nor NH₃ ionizes, and no OH⁻ is involved, so by the Arrhenius definition it isn’t a neutralization.
- Brønsted–Lowry. HCl gives H⁺ to NH₃, forming NH₄⁺ and Cl⁻ (which make up solid NH₄Cl). HCl is the proton donor, so it is the acid; NH₃ is the proton acceptor, so it is the base. It is an acid–base reaction.
- Conjugate pairs. HCl loses H⁺ to become Cl⁻, its conjugate base; NH₃ gains H⁺ to become NH₄⁺, its conjugate acid. The pairs are HCl/Cl⁻ and NH₄⁺/NH₃.
- Think about it. The Brønsted–Lowry definition covers more reactions: it explains this one, and the Arrhenius definition can’t.
Example 2: writing conjugates
Write the conjugate base of each acid: H₂SO₄, HCO₃⁻, NH₄⁺, H₂O. Then write the conjugate acid of each base: OH⁻, CO₃²⁻, H₂O, NH₃.
- The rule. Conjugate base = acid minus one H⁺ (charge goes down by 1). Conjugate acid = base plus one H⁺ (charge goes up by 1). Add or remove exactly one H⁺.
- Conjugate bases. H₂SO₄ → HSO₄⁻; HCO₃⁻ → CO₃²⁻; NH₄⁺ → NH₃; H₂O → OH⁻.
- Conjugate acids. OH⁻ → H₂O; CO₃²⁻ → HCO₃⁻; H₂O → H₃O⁺; NH₃ → NH₄⁺.
- Check the charges. HCO₃⁻ minus H⁺: (−1) − (+1) = −2, giving CO₃²⁻. NH₃ plus H⁺: 0 + (+1) = +1, giving NH₄⁺.
- Notice. H₂O and HCO₃⁻ appear in both lists: each has a conjugate base and a conjugate acid, so both are amphiprotic.
- A common mistake. Writing SO₄²⁻ as the conjugate base of H₂SO₄, which removes two H⁺ at once. SO₄²⁻ is the conjugate base of HSO₄⁻.
- HNO₃ in HNO₃ + H₂O → H₃O⁺ + NO₃⁻
HNO₃ gives H⁺ to water, so it is the acid.
- NO₃⁻ in the same reaction
NO₃⁻ is what is left after HNO₃ donates H⁺: the conjugate base of HNO₃.
- H₃O⁺ in the same reaction
H₃O⁺ is water after it has accepted H⁺: the conjugate acid of H₂O.
- H₂O in NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
Here water gives H⁺ to NH₃, so it is the acid.
- NH₃ in the same reaction
NH₃ accepts H⁺, so it is the base.
- OH⁻ in the same reaction
OH⁻ is what is left after water donates H⁺: the conjugate base of H₂O.
- OH⁻ in HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O
OH⁻ takes H⁺ from HCO₃⁻, so it is the base.
- HF in F⁻ + H₂O ⇌ HF + OH⁻
F⁻ accepts H⁺ and becomes HF, so HF is the conjugate acid of F⁻.
Common mistakes
Mistake 1: “A base must contain OH⁻.” That is the Arrhenius view. Under Brønsted–Lowry, anything that accepts H⁺ is a base: NH₃, CO₃²⁻, and F⁻ contain no OH⁻, but all are bases.
Mistake 2: “Any acid and any base in a reaction form a conjugate pair.” The two members of a conjugate pair differ by exactly one H⁺, like HCl and Cl⁻. HCl and OH⁻ differ by more than one H⁺, so they aren’t a conjugate pair.
Mistake 3: “Water is just the solvent, neither acid nor base.” Water accepts H⁺ from HCl, acting as a base, and gives H⁺ to NH₃, acting as an acid. Water is amphiprotic.
Quick check
In HSO₄⁻ + H₂O ⇌ SO₄²⁻ + H₃O⁺, what is the conjugate base of HSO₄⁻?
HSO₄⁻ loses one H⁺ to become SO₄²⁻. H₂SO₄ is the conjugate acid of HSO₄⁻, and H₃O⁺ is the conjugate acid of H₂O.
Which species can both donate and accept H⁺?
HCO₃⁻ can donate H⁺ to become CO₃²⁻ or accept H⁺ to become H₂CO₃. NH₄⁺ can only donate H⁺; Cl⁻ is an extremely weak proton acceptor; Na⁺ does neither.
Summary
- Arrhenius: acids produce H⁺ in water and bases produce OH⁻. The definition only covers water solutions.
- Brønsted–Lowry: acids donate H⁺ and bases accept H⁺. An acid–base reaction is a proton transfer, and H⁺ in water exists as H₃O⁺.
- A conjugate acid–base pair differs by one H⁺. Amphiprotic species such as H₂O and HCO₃⁻ act as acids in some reactions and bases in others.
On to AP (for later)
AP topic 8.1 defines acids and bases the Brønsted–Lowry way. You’ll be asked to write proton-transfer reactions, identify both conjugate pairs, and recognize that water is amphiprotic. AP questions usually write these equations with H₃O⁺. Topic 8.6 explains acid strength from molecular structure; lesson 2.12.2 starts with how far acids ionize.